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3.3 Ionic Bonding

GCSE Chemistry · AQA · Chapter 3: Structure and Bonding

A free GCSE Chemistry lesson for AQA (8462) from OBA School: electron transfer, the ionic bond and dot-and-cross diagrams for ionic compounds. Watch the video, read the explanation and worked examples, then try the revision questions. You can work through the full lesson, with every video, practice quizzes and printable worksheets with mark schemes, in our free online course.

What you will learn

  • Describe an ionic bond as the electrostatic attraction between oppositely charged ions
  • Explain the formation of ionic compounds in terms of electron transfer
  • Draw dot-and-cross diagrams for ionic compounds such as NaCl, MgO, CaCl₂, K₂O and LiF
  • Describe the limitations of dot-and-cross diagrams

Key words

  • Ionic bond: The strong electrostatic force of attraction between oppositely charged ions.
  • Electron transfer: The movement of electrons from a metal atom to a non-metal atom when an ionic compound forms.
  • Electrostatic force: The force of attraction between opposite charges (or repulsion between like charges).
  • Dot-and-cross diagram: A diagram that shows electrons from one atom as dots and from the other as crosses, so you can see where each electron came from.

What is an ionic bond?

Ionic bonding happens in compounds formed when a metal reacts with a non-metal. The metal atoms transfer their outer electrons to the non-metal atoms. This makes positive metal ions and negative non-metal ions.

Key fact: An ionic bond is the strong electrostatic force of attraction between oppositely charged ions. It acts in all directions.

The electrons are not shared. They move completely from one atom to another, so each ion ends up with a full outer shell.

Drawing dot-and-cross diagrams

Show the metal’s electrons as crosses (×) and the non-metal’s electrons as dots (•). Usually you only need to draw the outer shell. Follow these steps:

  • Work out how many electrons each metal atom loses and each non-metal atom gains.
  • Draw the non-metal ion with a full outer shell of 8: its own dots plus the crosses it has gained.
  • Draw the metal ion with its new outer shell (full) or with an empty outer shell. Both are accepted.
  • Put each ion in square brackets and write the charge at the top right.

Worked example: describing the formation of NaCl in words

A sodium atom loses its one outer electron to form a Na⁺ ion. A chlorine atom gains this electron to form a Cl− ion.
Both ions now have the electronic structure of a noble gas: Na⁺ is 2,8 (like neon) and Cl− is 2,8,8 (like argon).
The oppositely charged ions attract each other strongly.

Answer: transfer of one electron from Na to Cl; ions held by electrostatic attraction

Compounds with 2+ and 2− ions

When the charges are different sizes, you need more of one ion than the other. Count the electrons: every electron lost by a metal atom must be gained by a non-metal atom.

Use an electron bookkeeping table. The number of electrons lost by all the metal atoms must equal the number gained by all the non-metal atoms:

  • NaCl; Lost per metal atom 1; Gained per non-metal atom 1; Ratio metal : non-metal 1 : 1; Total transferred 1
  • MgO; Lost per metal atom 2; Gained per non-metal atom 2; Ratio metal : non-metal 1 : 1; Total transferred 2
  • CaCl₂; Lost per metal atom 2; Gained per non-metal atom 1; Ratio metal : non-metal 1 : 2; Total transferred 2
  • K₂O; Lost per metal atom 1; Gained per non-metal atom 2; Ratio metal : non-metal 2 : 1; Total transferred 2
  • Al₂O₃; Lost per metal atom 3; Gained per non-metal atom 2; Ratio metal : non-metal 2 : 3; Total transferred 6

Exam tip: Check the totals: in CaCl₂ calcium loses 2 electrons and the two chlorine atoms gain 1 each, so 2 lost = 2 gained.

Limitations of dot-and-cross diagrams

Dot-and-cross diagrams are good at showing where the electrons come from and the charges on the ions. But they:

  • do not show the 3D arrangement of ions in the giant lattice
  • do not show the relative sizes of the ions
  • suggest that the ions exist in separate pairs, when really each ion is surrounded by many oppositely charged ions
  • show electrons as dots and crosses, but all electrons are identical

Common mistakes

  • Drawing shared electrons between a metal and a non-metal. In ionic bonding electrons are transferred, not shared.
  • Forgetting the square brackets or the charges on the ions.
  • Drawing one oxygen atom taking one electron from each of two magnesium atoms. In MgO one Mg atom gives two electrons to one O atom.
  • Calling the ionic bond a “force between atoms”. It is between ions with opposite charges.
  • Saying the ions “want” a full shell. Say they achieve a full outer shell / noble gas structure.

Revision questions

Try these, then check your answers below each one.

Question 1. What is an ionic bond? (Choose from: The strong electrostatic attraction between oppositely charged ions, The attraction between positive ions and delocalised electrons, A shared pair of electrons between two atoms, A weak force between molecules.)

Answer: The strong electrostatic attraction between oppositely charged ions

Working: An ionic bond is the strong electrostatic force of attraction between oppositely charged ions. A shared pair is a covalent bond; ions in a sea of electrons is metallic bonding.

Question 2. When calcium chloride forms, how many electrons does each calcium atom transfer?

Answer: 2

Working: Calcium (2,8,8,2) loses 2 electrons, one to each of two chlorine atoms, forming Ca²⁺.

Question 3. In a dot-and-cross diagram of potassium oxide, K₂O, how many crosses (electrons that came from potassium) are in the outer shell of the oxide ion?

Answer: 2

Working: Two potassium atoms each give 1 electron, so the oxide ion has 6 dots and 2 crosses: 8 outer electrons.

Related lessons

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This lesson is section 3.3 (AGCCH0303) of our free GCSE Chemistry course for AQA. In the full course you get every video for this section, a self-marking practice quiz, a printable worksheet with a mark scheme, chapter tests, and your progress saved. Go to the full GCSE Chemistry course.

 

 

 

 

 

 

 

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